Lecture on Chemical Equilibrium

Introduction to Equilibrium

• Definition: Equilibrium in a chemical reaction occurs when the rate of the forward reaction equals the rate of the reverse reaction, resulting in constant concentrations of reactants and products.
• Dynamic Equilibrium: Even at equilibrium, reactions continue to occur with no net change in concentration of reactants and products.

Visualizing Equilibrium

• Example:
  • City A has 1000 cars and City B has 2000 cars, with a highway between them.
  • Each hour, 10 cars travel in both directions: A to B and B to A, representing equal rates in a reversible reaction.
  • Concentrations remain constant in both cities at equilibrium.

Concentration Profile

• Graphical Representation:
  • Concentration vs. Time graph for reactants (A) and products (B).
  • Initial conditions: 1 mole of A in a 1-liter container, no B present.
  • A decreases and B increases until both stabilize, indicating equilibrium.
  • Equilibrium is reached when changes in concentration stop.

Reaction Rates

• Graph of Reaction Rates:
  • Forward rate: k1 [A] decreases over time.
  • Reverse rate: k-1 [B] increases over time.
  • Equilibrium occurs when both rates are equal.

Equilibrium Constant (K)

• Derivation:
  • Forward reaction rate = k1[A], Reverse reaction rate = k-1[B].
  • At equilibrium: k1[A] = k-1[B], leading to K = k1/k-1 = [B]/[A].
  • K: Ratio of products over reactants, dependent on the reaction.

Types of Equilibrium Constants

• Kc (Concentration) and Kp (Pressure):
  • Kc: Equilibrium constant based on concentrations.
  • Kp: Equilibrium constant based on partial pressures.
  • Use in law of mass action for reactions.

Practice Problems

1. Write Equilibrium Expressions:

   • For a reaction A + B -> C + D, with coefficients j, k, l, m.
   • Kc = [C]^l[D]^m / [A]^j[B]^k.
   • Kp uses partial pressures instead of concentrations.

2. Calculate Kc Example:

   • Given N2 + 3Cl2 -> 2NCl3 with equilibrium concentrations.
   • Calculate Kc using concentration values.

3. Calculate Kp Example:

   • Given 2SO2 + O2 -> 2SO3 with equilibrium pressures.
   • Use Kp expression with given partial pressures.

4. Linking Kc and Kp:

   • Use the formula Kp = Kc(RT)^Δn.
   • Δn: Change in moles of gas (products - reactants).

5. Adjusting Reactions:

   • Doubling a reaction results in K squared.
   • Halving a reaction results in K^0.5.
   • Reversing a reaction results in 1/K.

6. ICE Table Method:

   • Use initial, change, and equilibrium (ICE) tables for determining equilibrium concentrations and K values.
   • Example: NOCl decomposing in a container to find Kc.

7. Complex Example Problems:

   • Given equilibrium conditions, calculate unknown concentrations or pressures using ICE tables and equilibrium expressions.
   • Different scenarios such as partial pressures and concentration changes are discussed.

8. Further Applications:

   • Calculating and interpreting K values when reaction conditions are altered (e.g., pressure, concentration, temperature changes).

Conclusion

• Emphasis on understanding equilibrium as a dynamic balance with constant activity and concentration stability in chemical reactions.
• Importance of equilibrium constants in predicting reaction outcomes and conditions.